Expansion of Octet. Why do elements do it?

Most of the elements do not have the noble gases' electronic configuration. Since the latter is generally unreactive, we have come to the conclusion that the octet configuration is stable and that the other elements would aim to achieve it. These elements can only achieve the octet configuration via chemical bonding; either forming of a covalent bond or an ionic bond with another atom.

However, we also noticed that there are compounds which clearly show that the element can form more bonds such that they have more than the 8 electrons around it. Compounds such as PCl₅ and H₂SO₄; where phosphorous and sulfur have 10 and 12 electrons around them respectively. Glaringly, all the elements which are found in the first and second period of the periodic table does not form compounds in which they have more than 8 electrons around them.

Therefore, these observations together with the knowledge of quantum shells, subshells and orbitals, we are able to come to the conclusion that elements found from Period 3 and beyond, have available empty orbital which must be energetically accessible to accept electrons for bonding. Thus, these elements are able to accept more electrons to bond with other elements. This results in the expansion of octet. In the case of PCl₅, phosphorous must have caused one of its valence electron to be excited to the 3d subshell (its energetically accessible orbital), so that it has 5 orbitals available for bonding and hence able to form a single bond each with 5 chlorine atoms.

However, is it that random, is it simply that elements which have energetically accessible orbital will exercise that privilege to use them? How do we account for compounds such as XeF₂?  Xe a clear noble gas element and yet it is willing to expand its octet, why would it do so? The expansion of the octet results in the elements to be able to make use of their energetically accessible empty orbitals. These orbitals are energetically slightly higher than the orbitals which the valence electrons are found. Hence, if the element wants to expand its octet, it will excite some of its valence electrons to the empty and energetically accessible orbital, so that they can form a bond with another atom via equal sharing of electrons or accept a pair of electrons from the donor to form a dative bond. If this is a favourable process, the energy from the formation of the bond must have compensated the energy needed to move a valence electron out of it the orbital that it resides in.

This is the complete reason to why the Period 1 and Period 2 elements do not expand their octet.  Their next empty orbital belongs to a different principal quantum shell. Hence, to make use of it for bonding would require a large amount of energy, and it is not compensated by the energy formed from the bond formation. Therefore, the combination of these two reasons allow us to observe expansion of octet in some elements and not all of them.

Interestingly, the cation of an ionic compound generally do not favour expansion of octet which results in formation of another neutral compound.  (For example adding more F₂ to MgF₂ to form MgF₄.) This is because it results in Mg to use more than 2 valence electrons for bonding.  This is energetically meaningless for the atom as it will have to invest in large amount of energy to make use of an inner shell electrons for bonding.  Hence, if cation wants to expand its octet, then it needs to have energetically accessible empty orbital to accept electron pairs from a donor.

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^{Article written by Kwok YL 2011.}

^{Disclaimer and remarks:}

• ^{If you would like to use this source, kindly drop me a note by leaving behind a comment with your name and institution. I am all for sharing as the materials on this blog is actually meant for the education purpose of my students.}
• ^{This material is entirely written by the author and my sincere thanks will be given to anyone who is kind, generous and gracious to point out any errors.}

Theory: Bond energy of A - B is the average of Bond energies of A-A and B-B

Let's try to make sense to one of the many theories in chemical bonding. We understand that the bond energy represents the amount of energy that is needed to break one mole of a covalent bond. Hence, the larger the bond energy, the stronger the bond. Interestingly, one of the theories in chemical bonding suggests that the bond energy of A-B is the average of the sum bond energy of A-A and the bond energy of B-B.

In this write-up, I will propose my interpretation to how this theory is obtained and hence using this as an angle to account for the anomalous situations. The Molecular Orbital (MO) theory best explains how the covalent bond is formed and this theory succinctly focuses on the overlapping of atomic orbital (Click here for some explanation to MO theory). While, at the GCE "A" level, we learn about overlapping of orbital for the formation of covalent bond; this is basically the fundamental and not the complete picture of MO theory. Nonetheless, this simple idea used in the "A" level is still competent to account for the theory of this topic.

Two assumptions are made when we talk about overlapping of atomic orbital to form covalent bond using the "A" level idea.  These will help to explain the theory in discussion:  

1. We assume that only a certain proportion of the atomic orbital will be used for bonding. Hence, when small atoms overlap with each other, the proportion of the overlap over the total area of the orbital is larger than for big atoms - that is where we obtain the term "effective overlap of atomic orbitals" to determine our strength of our covalent bonds. Strength of covalent bond is then told to us via bond energy. 
2. We assume that the atoms will always only make use of that same particular portion of their atomic orbital for bonding, regardless of the identity of the other atom.

These two propositions will suggest that if we half the bond energy of A-A, we can infer the area of the atomic orbital that the atom is using for bonding. This would be a fair suggestion, as a bond is made up of overlapping between two atomic orbital and the effectiveness the overlap (and hence the proportion used for overlapping) is told to us by the bond energy.

This explanation will nicely allow us to agree that the average of the sum of the bond energy of A-A and B-B will give us the bond energy of A-B.

However, when A and B is of different electronegativity, the electron cloud (which contains the bonded electrons) between A and B is pulled towards the atom which is more electronegative. You maybe incline to conclude that the area of the atomic orbital used for overlapping between A and B, may not be the same as A bonds with A and when B bonds with B (we can only ascertain this through theoretical calculations).  More importantly, the difference in electronegativity results in an added electrostatic attraction between partial charges that are formed in A and B; not accounted by the merely overlapping of atomic orbital.  The added electrostatic attraction creates our polar covalent bond.  A covalent bond with partial charges at each end.

Deviation to this theory is due to the weakness of the theory.  Essentially, it only accounts for a covalent bond formed between two atoms due to only the overlapping of atomic orbital. While, in the case of polar covalent bonds, we now know that there are two factors that result in this covalent bond to be formed:

1. Overlapping of atomic orbital and the constituent atoms; and 
2. Electrostatic attraction of partial charges due to the electronegativity difference of the constituent atoms.

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^{Article written by Kwok YL 2011.}

^{Disclaimer and remarks:}

• ^{If you would like to use this source, kindly drop me a note by leaving behind a comment with your name and institution. I am all for sharing as the materials on this blog is actually meant for the education purpose of my students.}
• ^{This material is entirely written by the author and my sincere thanks will be given to anyone who is kind, generous and gracious to point out any errors.}

Chemical Bonding - Strength of Interatomic Bonds (Part 2)

This article is an continuation of the article on interatomic bond. In this writeup, I shall discuss about the factors that affect the strength of the ionic bond and metallic bond. In trying to account for the strength of these bonds, it useful to make use of this rule: Always refer to the definition to how these bonds are formed first. (E.g. refer to the definition of covalent bond and then apply it in trying to account for the varying strength of different covalent bonds.)

(B) Strength of Ionic Bond.
Definition: Ionic bonds are formed because of electrostatic attraction between oppositely charge ions.

The following diagram illustrates how the ionic bond is affected. Strength of ionic bond is affected by lattice energy, which is directly affected by the product of the charges of the ions and inversely affected by the sum of the ionic radii.

(C) Strength of Metallic Bond. (using the sea of electrons model)
Definition: Metallic Bonds are formed because of electrostatic attraction between the cation of the metal and the sea of electrons.

Using the definition of metallic bonds, a metal atom which has more valence electrons will be able to contribute more electrons to the sea of the electrons. Hence, this result in a cation which has a greater positive charge. Therefore, the electrostatic attraction will be greater.

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Article written by Kwok YL 2009 (updated May 2009).}

^{Disclaimer and remarks:}

• ^{If you would like to use this source, kindly drop me a note by leaving behind a comment with your name and institution. I am all for sharing as the materials on this blog is actually meant for the education purpose of my students.}
• ^{This material is entirely written by the author and my sincere thanks will be given to anyone who is kind, generous and gracious to point out any errors.}

Chemical Bonding - Strength of interatomic bond (Part 1)

This article is an continuation of the article on interatomic bond. In this writeup, I shall discuss about the factors that affect the strength of the covalent bond. In trying to account for the strength of these bonds, it useful to make use of this rule: Always refer to the definition to how these bonds are formed first. (E.g. refer to the definition of covalent bond and then apply it in trying to account for the varying strength of different covalent bonds.)

(A) Strength of Covalent Bond:
Definition 1: Covalent bond is formed when there is electrostatic attraction between shared electrons and the nuclei of the two atoms.

When an atom is large, this will result that its valence electrons to be found further for the nucleus. In addition, the atom only make use of its valence electrons to form the covalent bond. Hence, when two larger atoms form a covalent bond with each other, their shared electrons will be significantly further away for the nucleus. Hence, the electrostatic attraction is weaker.

Definition 2: Covalent bond is formed when there is an overlapping between two atomic orbitals.

The atoms make use of their atomic orbital, where the valence electrons occupies, to overlap with each other and hence the a covalent bond is formed. The larger the atom, the less effective the overlap. Hence, the weaker the covalent bond.

Explanation - Strength of Covalent Bonds.
In addition, because of the definition of potential energy, this result in the covalent bond of H-H to be stronger and yet lower in potential energy than the Cl-Cl bond. This is succintly illustrated by the diagram below. Hence, lower potential energy implies greater electrostatic attraction between shared electrons and nuclei.

Lastly, bond energy refers to the amount of energy required to break a covalent bond. A larger bond energy implies that more energy is required to break a covalent. Hence, small atoms form strong covalent bond because the distance between shared electrons and nuclei is small. Therefore, greater electrostatic attraction, which results in more energy required to overcome this attraction, thus small atoms form bonds with large bond energies.

Other factors that affect strength of covalent bonds include: A polar covalent bond strengths the bond between two atoms. In addition, when two small atoms each having a lone pair of electrons, there is repulsion between the electron pairs hence, weakens the covalent bond.

In summary, using the other concepts you see in covalent bonding, the following flow-chart will be useful when we try to explain the different strengths of covalent bonds.


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Article written by Kwok YL 2009 (updated May 2009).}

^{Disclaimer and remarks:}

• ^{If you would like to use this source, kindly drop me a note by leaving behind a comment with your name and institution. I am all for sharing as the materials on this blog is actually meant for the education purpose of my students.}
• ^{This material is entirely written by the author and my sincere thanks will be given to anyone who is kind, generous and gracious to point out any errors.}

Chemical Bonding - Covalent bonds with a twist

Generally, the bond formed between a metal atom and a non-metal atom is that of an ionic bond. The metal atom will transfer its valence electron(s) to the non-metal atom such that both will become cation and anion respectively. In this manner, both will attain the favourable octet.

While, the bond between two non-metal atoms is that of a covalent bond. Neither atoms is willing transfer electrons to each other, hence to obtain the favourable octet both atoms would share a certain number of valence electrons. This sharing can take place when the atomic orbital of both atoms overlap.

(A) Seemingly ionic compound but are actually covalent.

AlCl₃ has covalent bonds between Al and Cl. This is contrary to what is expected.

If AlCl₃ was a ionic compound, the Al³⁺ cation is small and has a large charge, hence this cation is willing to polarise the electron cloud of the anion.

Cl^- is a large anion whose electron clouds are easily polarised. Polarisablity refers to the ease of distorting the electron cloud. When place next to Al³⁺, the electron cloud of Cl^- is distorted by Al³⁺. This severe distortion allows for both electron clouds to overlap. When that happens it mimics the overlapping of orbital hence covalent bond is formed.
Interesting, in a single molecule of AlCl₃, its dot and cross diagram does not fulfill the octet rule. This is because Al has only 3 valence electrons and even when all three are shared, there is only 6 electrons around Al. Hence, I must stress that this is not the norm. Usually, this rule is satisfied.

(B) Covalent bonds that carries some charge.

The assumption of a covalent bond is that the shared electrons are equally shared. This idea also suggests that the distance between the nucleus of one atom and the shared electrons is the same. However, this is not the case. Different elements' nucleus has different attractive power; some are more abled to pull electrons towards itself and hencce the term "electronegativity is introduced."

When the atom is more electronegative, it pulls the electrons closer to it. Hence when two atoms of differing electronegativity are bonded together, the atom which is more electronegative will pull the shared electrons closer to it.

This results in the less electronegative atom to carry a slight positive charge, while the other atom carries a slight negative charge. Since there is a small slight charges being produced,the two atoms are further held together because of electrostatic attraction between the partial charges.
In conclusion, we still need to bear in mind that the electrons are still shared between the two atoms; it is just now the shared electron is found closer to one nucleus than to the other. Therefore, a polar covalent bond is produced.

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Article written by Kwok YL 2009.}

^{Disclaimer and remarks:}

• ^{If you would like to use this source, kindly drop me a note by leaving behind a comment with your name and institution. I am all for sharing as the materials on this blog is actually meant for the education purpose of my students.}
• ^{This material is entirely written by the author and my sincere thanks will be given to anyone who is kind, generous and gracious to point out any errors.}

Chemical Bonding - Dative Covalent and Expansion of Octet

In these posts I shared about how interatomic bonds are formed and the factors affecting their strengths. While in this entry, I will be elaborating more about covalent bonds.

The earlier models of covalent bonds taught us that covalent bonds are formed when (1) the shared electrons used to formed the covalent bond must each come from the two constituent atoms and (2) the non-metals seek to attain an octet, hence search to have 4 electron pairs aroud it, therefore resulting in a maximum of 4 single bonds. These knowledge is still applicable, however, there are extensions.

(A) Dative covalent bonds.

Dative covalent bond are formed when an electron pair used to form a covalent bond between two particles come from just one source. Hence, this implies that one of the particles needs to have a pair of non-bonding electrons for donation and the other particle contains an empty orbital to accept the electron pair. I have illustrated using an example of an adduct formed by AlCl₃ and NH₃.

My second illustration of dative bond formed is using just AlCl₃. This compound usually exists as dimer and it is because Cl has available lone pair while Al has an empty orbital to accept the electrons. Interestingly, the Cl does not donate its lone pair of electrons to the Al it is attached to in a molecule; it rather donates its lone pair to a Al which is found on another molecule. This is because to form a double bond between Al and Cl is unfavourable as compared to having to form a dative covalent bond.

You may learn more about why AlCl₃ is a covalent compound over here.

(B) Expansion of octet - Having more than 8 electrons.

Elements found in Period 3 and below show an unique property. These elements show that the expansion of the octet is possible. For example, sulfur can exist as SO₂. In this molecule, S satisfies octet rule. While sulfur can also exist as SO₃ and in this molecule, S has more than 8 electron around it. This shows that S has expanded its octet (This is illustrated below.) It is because S contains empty low-lying 3d-orbtials which are available to accept more electrons.

Lastly, one needs to be mindful, when the octet is expanded, we will still get even number of electrons - all the electrons are paired. Hence, we will not observe 9 electrons or 11 electrons.
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Article written by Kwok YL 2009.}

^{Disclaimer and remarks:}

• ^{If you would like to use this source, kindly drop me a note by leaving behind a comment with your name and institution. I am all for sharing as the materials on this blog is actually meant for the education purpose of my students.}
• ^{This material is entirely written by the author and my sincere thanks will be given to anyone who is kind, generous and gracious to point out any errors.}

Chemical Bonding - Formation of multiple bonds

How are multiple bonds formed? We have learned that the number of electron pairs between two atoms indicate the number of electrons between the two atoms. It is easy to appreciate how a pair of electrons can be shared between two atoms. However, how is it possible for multiple pairs of electrons to be shared between two atoms?

Furthermore, when multiple bonds are formed, the strength of the each bond in the multiple bond is not identical. For example the C=O has two bonds between the C and O atoms; these two bonds are not equal in strength; one is weaker and the other is stronger.

Multiple bonds are formed when there is a sigma bond with the remaining bonds to be pi bonds. Hence, for example, a C=O contains one sigma bond and one pi bond. Do note that all single bond are sigma bond (the reason can be gathered from the later part of the post).

Sigma bonds are obtained when the orbital overlap in a head-on fashion. This can be illustrated by the diagram below.

While pi bonds are obtained when the orbital overlap in a side-on fashion. This type of bonding can also seen from the diagram below.

As the extend of orbital overlap in the sigma bond's case is greater than in the pi bond's case, hence it is not surprising that the sigma bond is stronger than the pi bond. This, therefore, accounts for the differing bond strength in the two bonds of the C=O.

An interesting fact, although this require some visualisation and hopefully this diagram helps: It is not possible to form a double/triple bond between two atoms if there is no sigma bond existing in the first place. Hence, essentially, all multiple bonds must have one sigma bond and the rest the bonds are pi.

In conclusion, in the "A" level curriculum, it is good to know that there are three bond types that we need to know: Single bond (contains only sigma bond), double bond (one sigma and one pi) and triple bond (one sigma and two pi).

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Article written by Kwok YL 2009.}

^{Disclaimer and remarks:}

• ^{If you would like to use this source, kindly drop me a note by leaving behind a comment with your name and institution. I am all for sharing as the materials on this blog is actually meant for the education purpose of my students.}
• ^{This material is entirely written by the author and my sincere thanks will be given to anyone who is kind, generous and gracious to point out any errors.}

Chemical Bonding - Interatomic Bonds

Chemical bonding is an important topic, the interatomic bonds helps to account for many physical properties, energies evolved or required for chemical reactions and these bonds also account for the speed of reactions. The three interatomic bonds which we will discuss are (1) Covalent bond, (2) Ionic Bond and (3) Metallic Bond.

(1a) Covalent Bond

Covalent bonds are formed when (A) there is electrostatic forces of attraction between shared electrons and the nuclei of the atom. However, in order for the two atoms to share electrons, the atomic orbtial of the constituting atoms must overlap with each other. Hence, (B) covalent bonds are formed when atomic orbitals overlap.

When two atomic orbital overlap, we obtained two molecular orbitals. (Since 1+1 = 2, we cannot just have one molecular orbital.) The two molecular orbitals are (i) Bonding orbital and (ii) Anti-bonding orbital. The illustration below provides greater details about the shape of (i) and (ii) and the relative energy level with respect to the atomic orbital.

However, if the the electrons are placed in the anti-bonding orbital, the bond breaks. (This is why in nucleophilic substitution, the nucleophile attacks the back of the C-X bond.)

P.S. The concept of Anti-bonding orbital is not in the GCE A level syllabus. The above picture is a basic illustration of Molecular Orbital Theory.

(1b) Multiple Covalent Bonds.

Using this diagram as our starting point, the number of electron pairs between two nuclei indicates the number of covalent bond between the two atoms. When there is one electron pair, this gives us a single bond. When there are two electron pairs between two nuclei, this gives us a double bond. Lastly, when there are three, it gives us a triple bond. Interestingly, we usually do not get more than 3 electron pairs between two nuclei. You can read more here.

(2) Ionic Bonds

Ionic bonds are formed because the electropostive atom has donated an electron to the electronegative atom. The former becomes a cation, while the latter becomes an anion. Electrostatic attraction between oppositely charged ions creates the ionic bond. Unlike covalent bonds, ionic bonds are not directional.

(3) Metallic Bonds

We are most commonly aware that metallic bond is formed because of electrostatic attraction between the array of cation and the sea of electrons. The illustration below depicts metallic bond and it helps in explaining the physical properties of metal.

However, when we boil liquid metal and form gaseous metal, do we obtain cations and electrons?

We don't. We actually obtain metal atoms, hence this is the limitation of the model which uses sea of electrons to describe metallic bonding. The more accurate illustration is given below but it is not in the A level syllabus. Do note that it is still a very simplified description.

Conclusion:

Electrostatic attraction results in chemical bonds to be formed, this attraction is a form of potential energy (PE). Due to the definition of PE, stronger the electrostatic attraction, the smaller the PE (note: numerical value gets larger).

As energy is conversed, formation of a bond (note: this refers to initially there is no bond between the two particles to a bond between the two particles.) results in a decrease in PE. Thus, there is a certain amount of energy lost when a chemical bond is formed which explains why bond formation is exothermic.

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Article written by Kwok YL 2009.}

^{Disclaimer and remarks:}

• ^{If you would like to use this source, kindly drop me a note by leaving behind a comment with your name and institution. I am all for sharing as the materials on this blog is actually meant for the education purpose of my students.}
• ^{This material is entirely written by the author and my sincere thanks will be given to anyone who is kind, generous and gracious to point out any errors.}